Transcript Chapter 4 - Aqueous Reactions

Chemistry 100
Aqueous Reactions
Solutions
A solution is a homogenous mixture of two
or more substances
 One substance (generally the one present
in the greatest amount) is called the
solvent
 The other substances - those that are
dissolved - are called the solutes

The Solution Process
Favourable interactions between the
solute and the solvent drive the formation
of a solution
 Example: NaCl (an ionic solid) dissolving
in water
 Water is a polar fluid (i.e., possesses a
permanent dipole)

Electrolytes
Salt is an ionic compound.
 NaCl is dissolved in water - the ions
separate.
 The resulting solution conducts electricity .
A solute with this property is called an
electrolyte

Strong Electrolytes
Strong electrolytes - completely
dissociated
 Some molecular compounds dissolve in
water to form ions.

 Dissolve
HCl (g) in water.
 All the molecules dissociate. So it is also a
strong electrolyte.
Weak and Nonelectrolytes
Weak electrolytes - only some of the
molecules dissociate, i.e., acetic acid
 Compounds that do not dissociate nonelectrolytes

 Sugars
 Ureas
 Alcohols
Acids
Acid - a substance that ionizes in water to
form hydrogen ions H+.
HCl (aq) H+ (aq) + Cl(aq)
 What is H+? A hydrogen atom without its
electron - a bare proton.

Monoprotic, diprotic, triprotic
One molecule of HCl gives one H+ ion:
HCl  H+ + Cl
We say that HCl is monoprotic - one
proton
 One molecule of sulphuric acid, H2SO4,
has two hydrogens to give away. It is
said to be diprotic.
 Phosphoric acid, H3PO4 is triprotic.

Some Chemical Structures
O
O
H
S
O
both H's ionize
O
H
H
C
H
O
C
OH
H
only this H ionizes
Acetic Acid

Generally write as CH3COOH, not
HC2H3O2.
 Weak
acid - doesn’t dissociate completely
CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq)
The double arrow - the system is in chemical
equilibrium!!!!
Bases
Bases are substances that accept (react
with) H+ ions. Hydroxide ions, OH, are
basic. They react with H+ ions to form
water:
H+ (aq) + OH (aq)  H2O (l)
 Ionic hydroxides like NaOH, KOH,
Ca(OH)2 are basic. When dissolved in
water they form hydroxide ions.

Ammonia solution
When ammonia gas dissolves in water,
some NH3 molecules react with water:
NH3(aq) + H2O(l) ⇄ NH4+ (aq) + OH– (aq)

NOTE - only some NH3 molecules
react with water. Ammonia is a
weak electrolyte.
Strong and Weak Acids and
Bases
Acids and bases that are strong
electrolytes are called strong acids and
strong bases.
 Strong acids are more reactive than weak
acids. Likewise for bases.
 Note exception - HF, a weak acid, is very
reactive

Acids you should know
Chloric acid
Hydrobromic acid
Hydrochloric acid
Hydroiodic acid
Nitric acid
Perchloric acid
Sulphuric acid
Acetic acid
HClO3
HBr
HCl
HI
HNO3
HClO4
H2SO4
CH3COOH (weak)
Bases you should know
Know the following bases:
Strong bases
a) Hydroxides of alkaline metals: LiOH,
NaOH, KOH
b) Hydroxides of the heavy alkaline earth
metals: Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak base: ammonia solution NH3

Metathesis reactions

A metathesis reaction is an aqueous
solution in which cations and anions
appear to exchange partners.
AX + BY  AY + BX
AgNO3 (aq)+ NaCl (aq)  AgCl (s) +
NaNO3 (aq)
Metathesis reactions (cont.)
Three driving forces
 Precipitate formation (insoluble
compound)
AgNO3(aq)+ NaCl(aq)  AgCl(s) +
NaNO3(aq)

Metathesis Reactions (Cont’d)
Weak electrolyte or nonelectrolyte
formation
HCl(aq) + NaOH(aq)  NaCl(aq) +
H2O(l)
 Gas formation
2HCl(aq) + Na2S(aq)  NaCl(aq) + H2S(g)

Neutralization
Mix solutions of acids and bases - a
neutralization reactions occurs.
acid + base  salt + water
 Salt does not necessarily mean
sodium chloride!!!!
 Salt - an ionic compound whose cation
(positive ion) comes from a base and
whose anion (negative ion) comes from
an acid

Precipitation Reactions
Some ionic compounds are insoluble in
water.
 If an insoluble compound is formed by
mixing two electrolyte solutions, a
precipitate results.

Precipitation (Cont’d)
Solubility - maximum amount of
substance that will dissolve in a specified
amount of solvent.
 Saturated solution of PbI2 contains 1 x 10-3
mol/L.
 A compound with a solubility of less than
0.01 mol/L - insoluble.
 More accurately - sparingly soluble.

Solubility Fact 1

All the common ionic compounds of the
alkali metals are soluble in water. The
same is true of the compounds containing
the ammonium ion, NH4+.
NaCl, K2CO3, (NH4)2S are all soluble
Solubility Fact 2
Salts containing the following anions are soluble
Anion
NO3
CH3COO 
Cl 
Br 
I
SO42
nitrate
acetate
chloride
bromide
iodide
sulphate
exception, salts of
none
none
Ag+, Hg22+,Pb2+
Ag+, Hg22+,Pb2+
Ag+, Hg22+,Pb2+
Ca2+, Sr2+, Ba2+,
Hg22+, Pb2+
Solubility Fact 3

Salts containing the following anions are
insoluble
Anion
S2
sulphide
CO32
carbonate
PO43
phosphate
OH 
hydroxide
exception, salts of
alkaline metal cations,
NH4+, Ca2+, Sr2+, Ba2+,
alkaline metal cations,
NH4+
alkaline metal cations,
NH4+
alkaline metal cations,
Ca2+, Sr2+, Ba2+,
Reaction forming gases
A metathesis reaction can occur due to the
formation of a gas which is not very
soluble in water.
 Examples involving hydrogen sulphide and
carbon dioxide

Reactions forming H2S
A metathesis reaction occurs when
hydrochloric acid is added to a sodium
sulphide solution.
2HCl(aq) + Na2S(aq)  H2S(g) + 2NaCl(aq)
 Net ionic reaction:
2H+(aq) + S2(aq)  H2S (g)

Reactions involving CO2

Carbonates and bicarbonates may be
thought of as the salts of carbonic acid
H2CO3 – unstable!!
H2CO3(aq)  CO2(g) + H2O(l)
Ionic Equations
Consider the reaction
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
 The above is known as the molecular
equation
 Note: the compounds are ionic (except
water)!!

Ionic Equations #2
Let’s show ionic compounds as ions
H+(aq) + Cl–(aq) + Na+(aq) + OH– (aq) 
Na+(aq) + Cl–(aq) + H2O(l)
 Some ions appear on both sides of the
equation.

Out with the spectators!
Remove ions that appear on both sides
H+ (aq) + Cl– (aq) + Na+ (aq) + OH– (aq) 
Na+ (aq) + Cl– (aq) + H2O (l)
 The unchanged ions are called spectators

The Net Ionic Equation

We are left with is the net ionic equation:
H+(aq) + OH– (aq)  H2O(l)
Note that the equation is balanced
for both mass and charge!!!
Another ionic reaction

Place zinc metal in a hydrochloric acid
solution – hydrogen is evolved!!
Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2 (g)
Why use ionic reactions?

They summarize many reactions.
 neutralization
of any strong acid by a
strong base is given by H+(aq) + OH– (aq) 
H2O(l)
The chemical behaviour of a strong
electrolyte  behaviour of its constituent
ions.
 Ionic equations can be written only for

Concentrations
How do we express the concentration of a
solution?
 Percentage is one way.

 2%
milk
 35% cream. (These are not true solutions)\
 Some beer is 5% alcohol

Note: % measurements can be %w/w,
%w/v, %v/v
Molarity
Must work in moles to do chemical
arithmetic.
 Chemists - molarity as their unit of
solution concentration

moles of solute
Molarity 
volume of solution (L)
Dilution

Dilute a solution
 more
solvent is added but the amount (mass
or moles) of solute is unchanged.
M1V1 = M2V2
The volumes can be either millilitres
(mL) or litres (L).
Ionic Concentration
NaCl in water - totally ionized into Na+ and
Cl ions.
 A 2.0 M NaCl solution

 Na+
concentration will be 2.0 M
 Cl concentration also 2.0 M

A 2.0 M solution of K2CO3,
 K+
concentration will be 4.0 M
 The concentration of CO32  2.0 M.
Oxidation and reduction
A piece of calcium metal exposed to the
air will react with the oxygen in the air
2Ca(s) + O2(g)  2 CaO(s)
 Ca has been converted to an ion Ca2+ by
losing two 2 electrons.
 Dissolve Ca in acid
Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g)
 Again the Ca has lost 2 electrons —
oxidation

Redox reactions
In the last two reactions, the Ca atom lost
two electrons. Where did they go?
 When one substance is oxidized, another
is reduced. An oxidation-reduction reaction
occurs. Or a redox reaction occurs.
 Oxidation: loss of electrons (more positive)
 Reduction: gain of electrons (less positive)

Oxidation of Metals - by air

Many metals react with oxygen in the air.
 Na


and K do so explosively!
Fe rusts - at a cost of $billions each year!
Aluminum oxidizes
 oxide
layer forms a skin which prevent further
oxidation. Al hides its reactivity.



Gold and platinum do not react with oxygen.
Silver tarnishes mainly because of H2S in
the air.
What does copper do?
Oxidation of Metals - by acids
Many metals react with acids:
metal + acid  salt + hydrogen gas
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
 Metals may also be oxidized by the
salts of other metals. Recall your lab
experiment
Fe(s) + CuSO4(aq)  Cu(s) +
FeSO4(aq)

Activity Series
We has seen that some metals react with
air, some also react with acids to give
hydrogen. We has also seen that some
metals can be oxidized by ions of other
metals.
 All this is summarized in the activity
series.

Activity Series
Li 
K
Ba 
Ca 
Na 
Mg 
Zn 
Fe 
Pb 
H
Cu 
Ag 
Li+
K+
Ba2+
Ca2+
Na+
Mg2+
Zn2+
Fe2+
Pb2+
 H+
Cu2+
Ag+
3+
+e
+e
+ 2e
+ 2e
+e
+ 2e
+ 2e
+ 2e
+ 2e




+e
+ 2e
+e
A metal can be oxidized
by any ion below it
Metals above H, react
with acids to give H2
The further up the
series, the more readily
the metal is oxidized
See your textbook
(p124) for more
elements
Some observations on the
series
Lead (Pb) is above H, so is Al. But
these metals are not attacked by 6M
HCl. They form very protective oxides.
 Cu reacts with nitric acid (HNO3)
because that acid is a strong oxidizing
agent in addition to being an acid.
 Gold (Au) and platinum (Pt) are
valuable because they are (a) rare and
(b) unreactive - they do not tarnish

Oxidation Numbers
Oxidation number - a fictitious charge
assigned to atoms either by themselves or
when combined in compounds as an
electron bookkeeping device.
 There are a number of simple rules that
chemists use to assign oxidation numbers.

Assigning Oxidation Numbers

In any elemental form (atom or molecule),
an atom is assigned a 0 oxidation number
 e.g.

He, Cu, N in N2, S in S8
For a monatomic ion, the oxidation
number equals the charge
 e.g.,
-1 for Cl in Cl-, +2 for Ca+2, -2 for S-2
Assigning Ox. Numbers (#2)

Fluorine’s oxidation number is -1 in any
compound.
 e.g.

-1 for F in CF4, but 0 for F in F2
Oxygen’s oxidation number is -2 except
when combined with fluorine or in
peroxides.
-2 for O in H2O and OH-, +2 for O in OF2,
-1 for O in H2O2
 e.g.
Assigning Ox. Numbers (#3)

For elements in Groups IA, IIA & most
of IIIA, oxidation numbers are positive
and equal to the group number.
 e.g.
+3 for Al in AlCl3, +1 for Na in NaCl, +2
for Mg in Mg SO4

Hydrogen has a +1 oxidation number.
Exceptions to this rule are the metallic
hydrides, in which it is -1.
 e.g.,
+1 for H in H2O and CH3OH, -1 for H
in NaH
Assigning Ox. Numbers (#4)

The sum of the oxidation numbers of the
atoms in a neutral compound is zero; in a
polyatomic ion, the sum equals the
charge.
 e.g.
see OH- and H2O above, +6 for S in SO4-2
Balancing Oxidation-Reduction
(Redox) Equations (#1)

Assign oxidation numbers to all atoms in
the equation.
 Note
- polyatomic ion that is unchanged in the
reaction may be treated as a single unit with
an oxidation number equal to its charge.
Balancing Redox Equations (#2)

Isolate the ATOMS that have undergone a
change of oxidation number
 A reduction
 An
in number indicates a reduction
increase in number, an oxidation
Balancing Redox Equations (#3)

Isolate the chemical species undergoing
oxidation/reduction (note: separate into an
oxidation and a reduction half-reaction).

Add the appropriate number of electrons
to the half-reactions
 Oxidation
 Reduction
– electrons on products side
– electrons on reactants side
Balancing Redox Equations (#4)

Remaining steps refer to the individual half
reactions
 Balance
for charges

Add H+ in acidic solution

Add OH- in basic solution
 Balance
water
the H and the O atoms by adding
Balancing Redox Equations (#5)

Balance the number of electrons in the
half-reactions
 Note:

electrons lost = electrons gained
Add the half-reactions, eliminating the
electrons and obtaining the complete
REDOX equation